Why is pH used as a measure of acidity if not only hydronium ions can act as a acid when in solution?

Question

The Brønsted-Lowry theory tells us that the when an acid such as $\ce{HCl}$ dissolves, the following reaction takes place: $$\ce{HCl + H2O <--> Cl- + H3O+}$$ It is said that the hydroxonium ion is what functions as an acid when a base is added to such a solution. However, there will indeed remain some $\ce{HCl}$ molecules in solution. So wouldn't they also act as an acid in some instances?

This problem becomes more apparent when we look at weak acids such as ethanoic acid. There are many many ethanoic acid molecules in comparison with the hydroxonium ions. Why is the acid 'weak' if the ethanoic acid molecules can just react with the base anyways?

To further explain where my problem is, consider the scale which we use to measure the 'acidity' of a solution, the pH scale. Consider a scenario where I have 2 different solutions of ethanoic acid and $\ce{HCl}$, both of the same concentration. The $\ce{HCl}$ solution will definitely have a lower pH and will be considered 'more acidic'. But would it really be correct to call the $\ce{HCl}$ solution 'more acidic' just because the $\ce{HCl}$ has dissociated to produce a much higher hydronium ion concentration? The ethanoic acid solution will have a lower hydrogen ion concentration, but the ethanoic acid molecules can still act as an acid … The main problem here is with using pH as a measure of acidity if molecules which haven't dissociated may also act as acids …

Answer 1

HCl being a strong acid, it is quantitatively dissociated in water, meaning that it fully dissociates into hydronium and chlorides. The hydronium are therefore the main contributors to the acidity of the solution. In a solution of ethanoic acid, the majority of acid molecules are undissociated. This is what "weak acid" means.

The amount of base necessary to neutralise a given acidic substance can be used as a scale of acid "strength" in chemical risk assessment.

Answer 2

It all boils down to the question which species can protonate which other species in a Brønsted-Lowry acid-base reaction. A lot of compounds have a generally somewhat acidic proton but it would be very wrong calling them an acid. For example, dihydrogenphosphate $\ce{H2PO3-}$ has a $\mathrm{p}K_\mathrm{a}$ value of $7.198$ — it can be deprotonated easily in aquaeous solution by sufficiently basic compounds but I really wouldn’t want to call it acidic. Or, to give another extreme example: furane can be deprotonated by n-butyllithium, but you really wouldn’t want to call it acidic, would you?

The good thing about the oxonium (or hydronium) ion concentration is that it is a pretty good approximation of the acidic effect. If I have a solution of acetic acid with $\mathrm{pH}\ 5$ and a (much more strongly diluted) solution of $\ce{HCl}$ with the same $\mathrm{pH}$-value, the protonating effect they have on different compounds will generally be the same.

And finally, $\mathrm{pH}$ is typically only used to measure the resulting acidity of a solution. When comparing the strengths of acids, $\mathrm{p}K_\mathrm{a}$ values are typically used. The picture being that with a $\mathrm{p}K_\mathrm{a}$-value difference of $\approx 3$, the weaker acid can be considered fully protonated at equimolar concentrations.

Answer 3

But would it really be correct to call the HCl solution 'more acidic' just because the HCl has dissociated to produce a much higher hydronium ion concentration?

Yes, it would be correct to call HCl more acidic because as an acid it is donating more protons compared to ethanoic acid now you argue that undissoaciated ethanoic acid can also react with a base to form salt. But salt formed will be of a weak acid and strong base(assuming), so this salt will hydrolyse and finally give back the acid again but this not the case if we use HCl , so no matter we have same concentration of different acids, more acidic will be one which gives more protons easily. So undissoaciated etanoic acid will remain again undisassociated.